CATALYSIS IN REDUCTION OF SILVER ION BY FERROUS ION
1503
CATALYSIS BY SILVER, BASIC CATALYSIS, AND CATALYSIS BY UNDISSOCIATED MOLECULES OF ACETIC ACID I N THE REDUCTION OF SILVER ION BY FERROUS ION BAL KRISHNA
AND
SATYESHWAR GHOSH
Physical Chemistry Laboratory, Allahabad University, Allahabad, India Received September $8, 1060
The reaction between silver ion and ferrous ion has been studied by Dhar (2) and by Roberts and Soper (16). Noyes and Brann (15) determined the equilibrium constant of the reaction: Fe++
+ Ag+ e Fe+++ + Ag
The present authors find that the reaction between silver ion and ferrous ion exhibits certain unique features from the point of view of the mechanism of general acid-base catalysis as conceived by Bronsted (I), and have therefore studied the reaction from various aspects. CATALYSIS BY SILVER AND THE ORDER OF THE REACTION
The catalytic role of silver has been reported by Sheppard (17). In a series of papers James (3, 4, 5 , 6, 7) has demonstrated that, in the reduction of silver salts, the reduced silver in the form of fine particles acts as a catalyst. Owing to a continuous flux in the quantity of the heterogeneous autocatalyst (wiz.,the h e particlesof silver) produced in the reaction, it becomes extremely difficult to determine the exact nature of the specific reaction rate constant. A direct determination of the velocity constant has been made by Krishna and Ghosh in the reaction between silver acetate and resorcinol and between silver tartrate and resorcinol (11, 12). In the reaction between silver ion and ferrous ion, the authors have already discussed the role of catalytic silver and the nature of the specific reaction rate constant (13, 14). I n the present paper they have studied the reaction with special reference to the role of hydrogen. ions, basic ions, and undissociated molecules of acetic acid. The effect of variation in the ionic strength of the medium upon the reaction velocity has also been studied. EXPERIMENTAL
Since the reduction of silver ion is a strongly photocatalytic reaction, the reaction was carried out in a bottle made of dark Jena glass. It is necessary to point this out, since others (2, 16) who have studied this reaction have neglected this precaution. Our conclusions are, therefore, applicable with certainty to the dark reaction only. The reaction is started by adding a known volume of standard ferrous sulfate solution to a buffered or nonbuffered solution of silver nitrate, silver sulfate, or silver acetate kept in the dark-glass bottle and maintained at a constant temperature. The buffer used is a mixture of sodium acetate and acetic acid. From this reaction mixture 10-ml. portions are taken out at definite intervals
1504
BAL KRISHNA AND SATYESHWAR QHOSH
of time and added to a known volume of a standard solution of potassium sulfocyanide. The latter removes silver ions and thus stops the reaction. From the supernatant filtrate 5 ml. is removed and added to 5 ml. of a standard silver nitrate solution. The excess of silver nitrate solution is then titrated against a standard solution of potassium thiocyanate in the presence of ferric alum as indicator. Titrations are carried out with a microburet. Table 1 shows the nature of the bimolecular velocity constant obtained in a typical case. A discussion of the exact nature of the specific reaction rate constant and of the order of the reaction is not relevant for the present purpose, particularly because these matters have been dealt with in previous papers (9, TABLE 1 The bimolecular velocity constant f o r the reduction of silver ion b y ferrous ion Reaction mixture: 0.02 M acetic.acid, 0.0025 M sodium borate, 0.005M ferrous sulfate, 0.005 M silver nitrate; temperature = 30°C.;pH = 7.00 TlYE
VOLUME OF
0.01 N
KCNS USID
mh.
ml.
0 10 20 30 50
2.10 2.08 2.00 1.96 1.80 1.68 1.59 1.53 0.52
80
120 180 Infinity
-
BIXOLECULAP VELOCITY CONSTANT
1 a(n
- x)
0.0009694 0.002124 0.002850 0.002944 0,002846 0.002522 0.001916
13, 14). The actual magnitude of the bimolecular velocity constant varies with time, owing to catalysis by silver. For the purpose of comparison, therefore, the maximum value of the velocity constant obtained in a given reaction has been taken as the measure of the velocity of the reaction. Table 1 shows that the value of the specific reaction rate constant increases with time, owing to catalysis by silver, which is a function of the primary particles of silver formed in the reaction. After reaching a maximum, the reaction constant begins to decrease, owing to the removal of primary particles as these coalesce to form secondary (11)and tertiary ones. DEPENDENCE OF REACTION RATE UPON ACETATE-ION CONCENTRATION
Table 2 contains a summary of the results obtained in the reaction between silver nitrate and ferrous sulfate in the presence of varying concentrations of acetate ion and a fixed amount of acetic acid. Similar results were obtained with silver sulfate and silver acetate. It might appear that the increase in velocity is due to an increase in hydroxylion concentration rather than acetate ions. I n order to make the velocity inde-
1505
CATALYSIS IN REDUCTION OF SILVER ION BY FERROUS ION
pendent of hydroxyl-ion concentration, the experiments were repeated with buffers of constant acid-base ratio but varying total concentration of the buffer. The results in table 3 xere obtained with silver acetate as the source of silver ions. This table proves conclusively that the reaction between silver ion and ferrous ion is catalyzed by acetate ions. In other words, it is a case of basic catalysis (1, 8), where anions of weak acids are involved in catalysis. TABLE 2 Reaction between silver nitrate and ferrous sulfate i n the presence of varying concentrations of acetate and a fixed amount of acetic acid 0.005 M silver nitrate: 0.005 M ferrous sulfate: 0.02 M acetic acid: temmrature = 30°C. CONCEXTPkIION OP ACETATE IOS6,
cA
,
!
YAXIYUY VALUE O F T H K B I Y O L E C W VELOCITY CONS~AN-I.
R
~
Y 0.05 0.10 0.15 0.20 0.30 0.40 0.60
CA
M 0.01
0.02 0.04
5.16 5.59 5.72 5.85 6.09
ij
R
0.01806 0.02450 0.03284 0.04230 0.06222 0.07722 0.15900
CA
I
R
M
0.03 0,05553 0.12990
0.023% 0.06887 0.11540
CATALYSIS BY BORATE AND CARBONATE IONS
If to the reaction mixture containing silver ions and ferrous ions, a crystal of sodium carbonate or borate is added, the reaction velocity is greatly increased.
No quantitative experiments (except in a few cases, see table l),however, can be carried out, owing to the high insolubility of silver borate and silver carbonate. It appears, therefore, that the reaction between silver ion and ferrous ion is a case of general basic catalysis. ANTICATALYFIC ACTIVITY OF HYDROGEN IONS
An increase in the concentration of hydrogen ions leads to a retardation in reaction velocity. This effect is observable if the concentration of acetic acid is increased in the absence of sodium acetate (see table 4). Similar results were obtained with sulfuric acid. Sulfuric acid (0.01 N ) com-
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BAL KRISHNA AND SATYESWAR QHOSH
CU.4
PH
K
M 0.02 0.04 0.08
4.83 2.95 2.94
0. ooo8641 0.0000000 (no reaction)
0.001023
TABLE 5 Effect of acetic acid upon the reaction velocity, when added in the presence of a constant amount of sodium acetate 0.005Msilver nitrate;0,005Mferrous sulfate; 0.05 Msodium acetate; temperature = 30°C. CU.4
M 0.02 0.05 0.08 0.12 0.22 0.42
I
DH
I
K
5.16 4.80 4.55 4.48 4.17 4.00
0.018060 0.003911 0.015050 0,03327 0.03327 0.059840
TABLE 6 Effect of variation i n the ionic strength upon lhe reaction velocity 0.005 M silver nitrate; 0.005 ill ferrous sulfate; 0.02 M acetic acid; 0.05 M sodium acetate; temperature = 30°C.
0.20
KK03
0.80
KNOs Mg(N0a)s
1.85
I I
-0.009410 -0.010476 -0.010730
EFFECT OF THE ADDITION OF ACETIC ACID IN THE PRESENCE OF SODIUM ACETATE; CATALYSIS BY UNDISSOCIATED MOLECULES OF ACETIC ACID
The fact of paramount importance in this connection is that acetic acid works as a negative catalyst only in the absence of sodium acetate, while in the presence of sodium acetate it acts as a positive catalyst. This we interpret to mean that the undissociated molecules of acetic acid also act as a positive catalyst. It will be observed in table 5 that there is at first a decrease and then an increase in the reaction velocity when acetic acid is added in the presence of a COP
CATALYSIS IN REDUCTION'OF SILVER ION BY FERROUS ION
1507
stant amount of sodium acetate. This presence of a minimum is of extreme importance in connection with catalysis by the undissociated molecule of acetic acid (see the discussion). At this point a number of objections may be raised. It may be argued that our assignment of a catalytic role to acetate ions and undissociated molecules of acetic acid is' only illusory, for an increase in the reaction velocity might well be due to a variation in the ionic strength of the medium or else might have been brought about by variation in the size and number of primary particles when different catalysts are added. Before we meet such objections, it is necessary to summarize the effect of variation in the ionic strength upon the reaction velocity when electrolytes are added (table 6). A K / A p is negative. This invalidates all arguments which ascribe a positive increment in reaction velocity to a salt effect. Therefore, the increase in velocity when the acetate-ion concentration is increased is due to its specific nature (Le., its basicity). It may be further pointed out that 10 ml. of M / 6 0 thorium nitrate in the reaction mixture completely stops the reaction. This is due to the fact that thorium ion behaves as an acid and its anticatalytic activity is like that of hydrogen ions. DISCUSSION
It is worthwhile to consider the mechanism of the salt effect in some detail. Bronsted (1) considers three kinds of salt effects: (1) a primary exponential salt effect, ( 2 ) a primary linear salt effect, and (3) a secondary salt effect. For our purpose we need not consider the primary linear salt effect, which is peculiar to reactions between neutral molecules or between a neutral molecule and an ion. The primary exponential salt effect is given by the formula (valid for water at 25°C.) h = KCnCB X 10zAzB"'
where h is the velocity of the reaction, C is the concentration of the reacting ions, 2 is the charge on the ions, and p is the ionic strength of the medium. I t is obvious that A h will be negative (if A p is positive) for ions of opposite charges, and positive for ions of similar charges. The equation is of great utility in tracing the reaction mechanism in ionic systems, and Krishna (10) has employed it in elucidating the mechanism of the reaction between ferric ion and stannous ion. I n the present case, both the reactants being positive ions, a positive salt effect should have been obtained. As a matter of fact, the reverse is the case. THE NEGATIVE SECONDARY SALT EFFECT
The real explanation of the negative salt effect observed in the present case lies in the fact that the dissociation of acetic acid increases in the presence of electrolytes (1) with a consequent increase in the concentration of hydrogen ions, which decrease the reaction velocity. The thermodynamic equilibrium constant for acetic acid is given by
1508
BAL KRISHNA AND SATYESHWAR GHOSH
where C is the concentration and f the activity coefficient. When g increases, the activity coefficients decrease and there is a consequent increase in the dissociation of acetic acid. The hydrogen-ion concentration therefore increases, and a decrease is observed in the reaction velocity. This negative salt effect has, perhaps, never been reported in the literature. On the addition of electrolytes the concentration of acetate ions is also increased, but the positive catalysis by acetate ions is more than balanced by the negative catalysis by hydrogen ions. Hence the negative salt effect is observed upon the addition of electrolytes. The second objection will now be considered. Do these external catalysts (viz., the acetate ions and the undissociated molecules of acetic acid) really act as homogeneous catalysts or do they play only a subsidiary role by altering the nature and extent of the catalytic surface? It is impossible to exclude the presence of silver from the reaction mixture, since silver is a reaction product. Consequently it is not possible to study separately the effect of homogeneous catalysts. Two points, however, appear to be decisive. Firstly, me find that in the presence of catalyzing agents (e.g., acetate ions) the induction period is considerably shortened and in fact can be indefinitely reduced by increasmg the quantity of the homogeneous catalyst. This lends support to the view that these factors work independently of silver, for an induction period means a period of very slow reaction in almost complete absence of silver particles. An induction period therefore implies an interval of homogeneous reaction when the catalytic surface is negligibly mall, and a shortening of the induction period means that the catalyzing agents (other than silver) are working independently of silver. Secondly, we find that on the addition of acetic acid in the presence of a constant amount of sodium acetate, at first there is a decrease (see table 5) in the value of K , and after a minimum is reached, there is again an increase. Further work is being carried out in this laboratory in this connection (to be published later), and we find that there is always a minimum point in the value of K when acetic acid is added in the presence of a constant amount of sodium acetate. It is, a priori, improbable that the catalytic surface should at first decrease and then increase as the amount of acetic acid is increased. On the other hand, our theory of homogeneous reaction explains it mathematically as follows: The overall bimolecular velocity constant can be split up into two termsone heterogeneous and the other homogeneous-and we can write :
K =
Khorno.
+
(1)
Khetero.
In the presence of acetic acid and sodium acetate the K h o m o . itself is made up of three terms-one due to the catalytic activity of undissociated molecules of acetic acid, the second due to that of acetate ions, and the third due to hydrogen ions (this last term being negative). We can, therefore, write Khomo.
= KI(1
- a)C + KI(CA
a C ) , f KaaC
(2)
where a is the degree of dissociation of acetic acid, (1 - a) is the fraction of the undissociated molecules of acetic acid, CA is the concentration of acetate ions
CATALYSIS I N REDUCTION O F SILVER ION BY FERROUS 10s
1509
(this being constant), and C is the concentration of acetic acid. It is to be noted that the value of Ka is negative, since hydrogen ions act as an anticatalyst. Writing the ordinary dissociation constant of acetic acid in the presence of sodium acetate we have:
+
K = aC(aC C,) C(1 - a) or C=
K(l
-
a)
- aCA
a:
Substituting the value of C in equation 2 we have
Khomo. =
- UCA]4-K z [ K ( l - a) - aCAl
K 1 ( l - a) [ K ( 1- a) d
U
3-
Ka[K(1 - a)
- aCd + K 2 C A
U
= +(a), say.
(3)
The first condition for a minimum for Khomo.is that
Differentiating equation 3 and equating it to zero we have -=--
da
2K1K a
3
2K1K +
7
+
-a2- - -K3K -
KLCA- KtK 7
-
whence
2K1 K a = 2K1K
- K z K - KaK
+ KICA
Now since the anticatalytic activity of hydrogen ions is much greater than the positive catalytic activity of acetate ions (0.01 N sulfuric acid completely stops the reaction in the absence of sodium acetate), we have and since K3 is negative, - K X
>
0, and
IKnXI /
